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Mastering Group 1: Alkali Metals and Their Properties
Dive into the fascinating world of alkali metals! Discover their unique properties, reactivity patterns, and applications in chemistry. Learn how these Group 1 elements shape our understanding of the periodic table.

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Now Playing:Group 1 alkali metals – Example 0a
Intros
  1. Group 1 and group 2: a summary
  2. Group 1 and group 2: a summary
    Intro to group 1 and 2.
  3. Group 1 and group 2: a summary
    Properties of group 1 and group 2 metals.
Examples
  1. Recall the trend in properties of the alkali metals.
    Sort the alkali metals for the following properties:
    i) From hardest to softest.
    ii) From lowest to highest melting point.
    Group 1 and group 2: Alkali and alkaline earth metals
    Jump to:NotesConceptFAQsPrerequisites
    Notes
    In this lesson, we will learn:
    • To recall properties of the alkali metals
    • To understand the trend in properties found in the alkali metals
    • To apply knowledge of electronic structure and bonding to explain the trends in group 1 properties.
    • To know the common reactions of group 1 and group 2 elements.

    Notes:
    • We saw earlier that the Periodic Table is arranged, left to right, by proton number and number of outer shell electrons. The number of outer shell electrons dictates the chemical properties of an element.
      Therefore, it is easy to see which elements have similar properties to each other – they will be in the same column of the table as each other, the columns which we call groups.
      Groups 1 and 2 have 1 and 2 electrons in the outer shell, respectively. Their behaviour in chemical reactions is similar because of the similar outer shell configuration.

    • The alkali metals in group 1 and the alkaline earth metals in group 2 are very well-studied groups of elements, with clear patterns in how their properties change.
      They have been known for quite a long time through their compounds, like their metal oxides, because they are reactive and quickly form compounds with oxygen in air. Their oxides all produce alkaline solutions in water, which is how they get their names. Because of their reactivity, they were not isolated as metals until later on.

    • Alkali metals and alkaline earth metals have the following properties:
      • They are relatively soft metals.
      • They are relatively low density metals.
      • They have relatively low melting points compared to metals in general.
      • They are reactive, more so than d-block metals in general and they react vigorously with water.
        • The products of this reaction are hydrogen gas and a metal hydroxide – this forms an alkaline solution, which gives the two groups their names.

    • As you go down the group, the properties of the elements change in the following ways:
      • The metals get softer.
      • The melting point of the metals gets lower.
      • The metals get denser.
      • The metals get more reactive.
      • Ionization energies of the elements decrease.

    • The alkali metals all have a valence of 1 and alkaline earth metals have a valence of 2. As they are metals, they form ionic compounds with non-metals. In these compounds, you’ll see alkali metals with a 1+ charge (such as in NaCl), whereas alkaline earth metals will hold a 2+ charge (such as in MgCl2).

    • As you descend the table, a similar trend in the change in properties is observed in both groups. For example, both group 1 and group 2 show a decrease in ionization energies going down the group.

      As explained in Periodic Trends: Ionization Energy, this is due to greater shielding from more inner electron shells between the nucleus and the outer shell. Greater shielding makes losing the one (in group 1) or two (in group 2) outer shell electrons increasingly easy, and therefore reactivity in general increases going down the two groups.

    • The reactions of group 1 and group 2 metals are very similar because they have similar outer shell configurations and elements in both groups are prone to losing their outer shell electrons.
      • One of the most common reactions with group 1 and 2 elements is the reaction with water. Group 1 and group 2 metals are both known for reacting vigorously with water compared to other metals. The equation can be written generally as:

        For a group 1 metal:
        2M (s) + 2H2O (l)  \, →  \, 2MOH (aq) + H2 (g)


        For a group 2 metal:
        M (s) + 2H2O (l)  \, →  \, M(OH)2 (aq) + H2

        The stoichiometry is slightly different because two hydroxide groups will be bonded to the group 2 metal because it forms a 2+ ion, unlike the 1+ ion of a group 1 metal.


      • Another common reaction is the group 1 or 2 reaction with oxygen. This can be written generally as:

        For a group 1 metal:
        4M (s) + O2 (g)  \, →  \, 2M2O (s)


        For a group 2 metal:
        2M (s) + O2 (g)  \, →  \, 2M2O (s)

        • Again, the stoichiometry is slightly different in the two reactions.
        In the reaction with heavier group two metals, a metal peroxide can form, such as with barium and strontium:
        M (s) + O2 (g)  \, →  \, MO2 (s)

        The peroxide ion (-O-O-) has two oxygen atoms, each with a weak covalent bond to the other, and a 1- charge (O22- overall). It does not form in this reaction with the smaller group 2 metals such as Be or Mg because they form highly polarizing ions.
        These are small with a very high charge density which draws the 2- negative charge in the peroxide ion towards it with great strength (they polarize the ion). The peroxide ion becomes a stable O2- oxide ion and the weak O-O covalent bond breaks.
        The heavier, less dense and less polarizing ions like Sr2+ and Ba2+ cannot do this, so the metal peroxide is stable.

      • The reaction with oxygen to form a metal oxide occurs spontaneously in air and when the metals are heated by flame, showing a distinct colour.

        • The following is a table of the flame colours observed in the group 1 and 2 elements:

        Group 1 elements

        Group 2 elements

        Lithium: red

        Beryllium: white/colourless

        Sodium: orange/yellow

        Magnesium: white/colourless

        Potassium: lilac

        Calcium: red/orange

        Rubidium: red

        Strontium: dark red

        Caesium: blue/violet

        Barium: green


      • The reaction between the group 1 and 2 metals with chlorine can be written generally as:


        • For a group 1 metal:
        2M (s) + Cl2 (g)  \, →  \, 2MCl (s)


        For a group 2 metal:
        M (s) + Cl2 (g)  \, →  \, 2MCl (s)


    • The solubility of many group 1 and 2 metal compounds have trends down the groups. Generally, group 1 metal compounds are more soluble than any group 2 analogue.
      For example, you would predict that potassium hydroxide has greater solubility (in water at a fixed temperature) than calcium hydroxide.
      • Solubility of the metal hydroxides increases down the group. This is true for both group 1 and group 2 hydroxides, and as said above the group 1 hydroxides are much more soluble
      • Solubility of the group 2 metal sulfates decreases down the group. Barium sulfate is the product we look for in the common test for sulfate ions. The white precipitate is obvious as BaSO4 is extremely insoluble.
      • Solubility of the group 2 metal carbonates generally decreases down the group. In group 1, solubility actually increases down the group.
      Explaining the solubility trends is not necessary. In short, it’s because of two enthalpy measurements (lattice enthalpy and enthalpy of hydration) that change at different rates down the groups, depending on the different anions (sulfate, carbonate, etc) used... don’t worry about it!
    Concept

    Introduction

    Alkali metals and alkaline earth metals, found in Groups 1 and 2 of the periodic table elements, respectively, are essential elements in chemistry. The introduction video provides a comprehensive overview of these elements, offering valuable insights into their unique properties and behaviors. Alkali metals, including lithium, sodium, and potassium, are highly reactive and soft, while alkaline earth metals, such as beryllium, magnesium, and calcium, exhibit similar but less extreme characteristics. These elements play crucial roles in various chemical reactions and have numerous applications in industry and everyday life. Their distinctive properties, such as low ionization energies and high reactivity with water, make them fascinating subjects of study. Understanding the behavior of alkali and alkaline earth metals is fundamental to grasping broader concepts in chemistry and the periodic table elements. The video serves as an excellent starting point for exploring these important elements and their significance in the field of chemistry.

    FAQs

    Here are some frequently asked questions about alkali and alkaline earth metals:

    What are the 5 properties of alkaline earth metals?

    The five key properties of alkaline earth metals are: 1) They are silvery-white in color, 2) They have relatively low densities, 3) They are good conductors of heat and electricity, 4) They have relatively high melting and boiling points compared to alkali metals, and 5) They are highly reactive, though less so than alkali metals.

    Are alkaline earth metals soft or hard?

    Alkaline earth metals are generally softer than most other metals but harder than alkali metals. Their hardness increases as you move down the group in the periodic table. For example, beryllium is relatively hard, while barium is quite soft.

    What are the main characteristics of alkali metals?

    The main characteristics of alkali metals include: 1) They are extremely reactive, 2) They have low melting and boiling points, 3) They are soft and can be cut with a knife, 4) They have low densities and can float on water, 5) They have one valence electron which they readily lose in chemical reactions.

    What is the difference between alkali and alkaline earth metals?

    The main differences are: 1) Alkali metals are in Group 1 of the periodic table, while alkaline earth metals are in Group 2, 2) Alkali metals have one valence electron, while alkaline earth metals have two, 3) Alkali metals are more reactive than alkaline earth metals, 4) Alkali metals form +1 ions, while alkaline earth metals form +2 ions, 5) Alkali metals are softer and have lower melting points than alkaline earth metals.

    Why do alkaline earth metals have similar chemical properties?

    Alkaline earth metals have similar chemical properties because they all have two valence electrons in their outermost shell. This electronic configuration leads to similar behavior in chemical reactions, such as their tendency to form +2 ions and their reactivity with water and oxygen. Additionally, as you move down the group, the atomic radius increases, but the effect of the two valence electrons remains consistent, resulting in gradual trends in properties.

    Prerequisites

    Understanding the foundations of chemistry is crucial when delving into the fascinating world of Group 1 and Group 2 elements, also known as alkali and alkaline earth metals. To fully grasp the properties and behaviors of these important element groups, it's essential to have a solid understanding of two key prerequisite topics: the history and development of the periodic table and atomic structure.

    The periodic table is the cornerstone of modern chemistry, and its evolution provides invaluable context for understanding how elements are organized. By exploring the history and development of the periodic table, students gain insight into how scientists discovered patterns in elemental properties, leading to the grouping of elements with similar characteristics. This knowledge is particularly relevant to Group 1 and Group 2 elements, as their placement in the periodic table reflects their shared properties and reactivity trends.

    Moreover, comprehending the atomic structure is fundamental to understanding why alkali and alkaline earth metals behave the way they do. The electronic configuration of these elements, particularly their valence electrons, plays a crucial role in determining their chemical and physical properties. By mastering concepts related to atomic structure and bonding, students can better explain phenomena such as the high reactivity of Group 1 metals or the tendency of Group 2 elements to form divalent compounds.

    The periodic table groups are not arbitrary divisions; they are the result of careful observation and analysis of elemental properties. Understanding how these groups were established helps students appreciate the logic behind the classification of alkali and alkaline earth metals. It also provides context for why these elements exhibit gradual changes in properties as you move down each group.

    Additionally, a strong foundation in atomic structure and bonding is essential for grasping concepts like ionization energy, atomic radius, and electronegativity all of which are critical in explaining the characteristic behaviors of Group 1 and Group 2 elements. For instance, the low ionization energies of alkali metals can be directly linked to their atomic structure, explaining their high reactivity and tendency to form positive ions.

    By thoroughly understanding these prerequisite topics, students will be well-equipped to explore the intricacies of alkali and alkaline earth metals. They will be able to draw connections between the elements' positions in the periodic table, their electronic structures, and their observed chemical and physical properties. This comprehensive approach not only enhances learning about Group 1 and Group 2 elements but also provides a solid framework for understanding broader concepts in chemistry.